Examples:
ΔH ≈ -814 kJ/mol (exothermic)
- Energy to break bonds = 4 × 413 + 2 × 498 =
2,648 kJ - Energy released forming bonds = 2 × 805 + 4 × 463 =
3,462 kJ - ΔH = broken − formed = 2,648 − 3,462 =
-814 kJ/mol
Average bond enthalpies are for the gas phase, so the answer is an estimate — typically within a few per cent of measured values for gas-phase reactions.
How it works
Breaking bonds takes energy and forming bonds releases it, so ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Draw displayed formulas first so you can count every bond; a double bond counts once, using its own C=C or C=O value.
Average bond enthalpies are averaged over many molecules and apply to the gas phase, so the result is an estimate. For methane combustion it gives about −814 kJ/mol against a measured −802 kJ/mol for gaseous water (−890 kJ/mol if the water condenses). For exact values, use enthalpies of formation.
Frequently asked questions
- Why is it "bonds broken minus bonds formed"?
- Breaking bonds absorbs energy (positive) and forming bonds releases energy (negative). If more energy is released than absorbed, ΔH is negative and the reaction is exothermic.
- Why doesn't my answer match the data book?
- Bond enthalpies are averages, and they ignore changes of state. Reactions involving liquids or solids, such as water forming as a liquid, differ most.
- Which C=O value should I use?
- 805 kJ/mol for the C=O bonds in carbon dioxide, and about 745 kJ/mol for C=O in aldehydes and ketones. Exam boards may give their own values; use those if provided.
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